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Chapter 2 Section 3 Bohr Atom and Electron Configuration Notes
Name: ________________________________
Date: _______________ Pd.: _____
Introduction
An electron’s position in an atom or ion can be described by determining its electron configuration and orbital diagram.
These representations of an atom or ion can explain physical and chemical properties of the substance, including
magnetic attraction.
Bohr Atom
Niels Bohr’s “Planetary Model”- in which the nucleus is surrounded by orbitals of electrons - resembles the solar system.
 Electrons could be excited by energy and move to an outer orbit (excited level).
 They could also emit (give off) radiation when falling to their original orbit (ground state).
 Basic Components of the Bohr Model:
o Energy Levels: Energy levels are the volume of space where certain electrons of specific energy are restricted to
move around the nucleus.
 Energy levels are categorized by the letter n (lowercase n) using whole numbers. (n=1, 2, 3, etc.)
o Orbitals: An orbital is the space where one or two paired electrons can be located or the probability of an
electron’s location.
 Orbitals are represented by letters (s, p, d, f, etc.) (lowercase letters)
o Outer or Valence Shell: The valence shell is the last energy level containing loosely held electrons. These are the
electrons that engage into bonding and are therefore characteristic of the element’s chemical property.
(You do not need to memorize the shapes.)
Electron Configuration
Electron Configurations describe the exact arrangement of electrons (given as superscripts) in successive energy levels
or shells (1, 2, 3, etc.) and orbitals (s, p, d, f) of an atom, starting with the innermost electrons.
o Example: A lithium atom’s configuration is 1s22s1
o Superscripts mean two electrons are in the 1s orbital and one electron is in the 2s orbital.
Several Rules are applied to the filling of electrons:
o Pauli Exclusion Principle: The Pauli Exclusion Principle states that an orbital can hold a maximum of two
electrons is they are of opposite spins. In other words, every electron has a unique set of quantum numbers.
o Hund’s Rule: Hund’s Rule states that the most stable arrangement of electrons in the same energy level in which
electrons have parallel spins (same orientation).
o Aufbau Principle: The Aufbau Principle is based on the Pauli Exclusion Principle and states that electrons are
placed in the most stable orbital.
 Aufbau means “building up” in German.
 Example: 1s2, 2s2, 2p6
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How is an Electron Configuration written?
o Determine the number of total electrons.
o This is the same as the Atomic Number for neutral atoms.
o For ions (charged atoms), the total electron is corrected for the charge.
 Ions are atoms or molecules which have gained or lost one or more valence electrons, giving the
ion a positive or negative charge.
 For anions (negative ions) – add electrons
aNion (a Negative ion)
 For cations (positive ions) – subtract electrons
ca+ion (positive sign! = positive ion)
o The electrons are “added” according to Hund’s Rule and the Aufbau Principle. (see next diagram)
o Keep in mind the maximum number of electrons in each type of orbital
 S subshell has 1 orbital & holds 2 electrons
s ____
2 electrons per orbital (line)
 P subshell has 3 orbitals & holds 6 electrons
p ____ ____ ____
 D subshell has 5 orbitals & holds 10 electrons
d ____ ____ ____ ____ ____
 F subshell has 7 orbitals & holds 14 electrons
f ____ ____ ____ ____ ____ ____ ____
__________________________________________________________________________________________________
1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p8s
Draw the above on a piece of paper -the chart and the order written out. Put in your binder after this page of notes.
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Examples: Give the electron configuration for Na, Ta, O2-, and N2+
N: (Total electrons – 7)
1s ___ 2s ___ 2p ___ ___ ___
Ta: (Total electrons – 73)
1s ___ 2s ___ 2p ___ ___ ___
2O : (Total electrons – 8 + 2e-= 10 electrons)
1s ___ 2s ___ 2p ___ ___ ___
1+
Na : (Total electrons – 11 - 1e- = 10 electrons
1s ___ 2s ___ 2p ___ ___ ___
Configuration: 1s22s22p3
Configuration:
Configuration:
Configuration:
 Notice that O2- and Na1+ have the same number of electrons. These ions are isoelectric (the same number of
electrons)
 Notice that in Nitrogen that the 2p orbital is not filled all the way up. It has 3 electrons when it could potentially hold
6. These extra electrons are called valence electrons (Electrons added since the last noble gas). They determine how
reactive an atom is. The closer an atom is to having a full amount of electrons in its outer orbital, the more reactive
the atom is. Once the electrons completely fill that outer filled noble gas orbital, they are inert (non-reactive) and
stable.
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Condensed Electron Configuration (Shorthand Electron Configuration)
Figure 10.1 shows the electron configurations for the valence electrons only (*Condensed).
Shorthand Electron Configuration Rules:
1. Find the element on the periodic table.
2. Go backward until you reach the first smaller noble gas.
3. Put that element (noble gas) in brackets ( Helium = [He] ). Complete the electron configuration from this point.
Orbital Diagrams
Now that you understand electron configurations, an orbital diagram can be drawn. Orbital diagrams represent the
orbital where each electron is located. An arrow is used to represent each electron spinning in a particular direction.
Recall that s subshells have 1 orbital, p subshells have 3 orbitals, and d subshells have 5 orbitals. The orbital diagram can
be constructed using the electron configuration and understanding Hund’s Rule, stating each orbital of a subshell is
filled first with an unpaired electron and then paired.
Aluminum Example of Orbital Diagram:
These are two examples of the way
to imagine the electron
configuration orbital diagrams.

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Magnetism
Particular elements show properties of magnetism when
placed in a magnetic field. Paramagnetic elements are
elements that are attracted to a magnet, and paramagnetism
is connected with unpaired valence electrons. Notice Iron (Fe)
in the above diagram has four unpaired valence electrons.
This indicated that iron will strongly attract to a magnetic
field. In contrast, Diamagnetic elements are elements that
have all paired electrons and do not possess the ability to
attract to a magnetic field. Examples of diamagnetic elements
are the noble gases and O2-.
Quantum Numbers
Principal Quantum Number
Orbital Angular Momentum
Quantum Number
Magnetic Quantum Number
Magnetic Spin Quantum
Number
n
l
(cursive lowercase L)
ml
ms
The number before the letter
The orbital (s, p, d, f) but in a
number form that starts from
zero.
The maximum number under
the orbital lines
2s ___
0
2p ___ ___ ___
-1 0 1
3d ___ ___ ___ ___ ___
-2 -1 0 1 2
Electron Spin
(Was the last arrow you drew
pointing up or down?)
Lithium 1s22s1
s=0
p=1
n=2
d=2
f=3
ml = 0 then, l = 0
ml = 1 then, l = -1, 0, 1
ml = 2 then, l = -2, -1, 0, 1, 2
ml = 3 then,
l = -3, -2, -1, 0, 1, 2, 3
Spin up with ms = +1/2 or
Spin down with ms = −1/2
The principal quantum number, n, designates the principal
electron shell. Because n describes the most probable
distance of the electrons from the nucleus, the larger the
number n is, the farther the electrons are from the nucleus,
the larger the size of the orbital, and the larger the atom is.
The orbital angular momentum quantum number, l,
determines the shape of an orbital, and therefore the angular
distribution.
The magnetic quantum number determines the number of
orbitals and their orientation within a subshell.
The electron spin quantum number ms designates the
direction of the electron spin and may have a spin of +1/2,
represented by↑, or –1/2, represented by ↓. This means that
when ms is positive the electron has an upward spin, which
can be referred to as "spin up." When it is negative, the
electron has a downward spin, so it is "spin down." The
significance of the electron spin quantum number is its
determination of an atom's ability to generate a magnetic
field or not.
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